THE  OXIDATION  OF  FERROUS  SULPHATE  TO 
FERRIC  SULPHATE  BY  MEANS  OF  AIR 


By 

JAMES  STEWART  MACHIN 

A.  B.  Westminster  College,  1921 


THESIS 

SUBMITTED  IN  PARTIAL  FULFILLMENT  OF  THE  REQUIREMENTS 
FOR  THE  DEGREE  OF  MASTER  OF  SCIENCE  IN  CHEMISTRY 
IN  THE  GRADUATE  SCHOOL  OF  THE  UNIVERSITY 
OF  ILLINOIS,  1922 


URBANA,  ILLINOIS 


Digitized  by  the  Internet  Archive 
in  2015 


https://archive.org/details/oxidationofferroOOmach 


I 922 

M i a 


o. 

■<r 


UNIVERSITY  OF  ILLINOIS 


s\S 

a. 

01 


THE  GRADUATE  SCHOOL 


JULY  28, 


192 


I HEREBY  RECOMMEND  THAT  THE  THESIS  PREPARED  UNDER  MY 


SUPERVISION  BY 


JAMES  S-T-S^ART -MAORI  N 


ENTITLED. 


THE  OXIUAT I OH.  GY 


iS-  SULPHATE-  TO 


FERRIC 


MEA-K3  QF-AIR 


BE  ACCEPTED  AS  FULFILLING  THIS  PART  OF  THE  REQUIREMENTS  FOR 


THE  DEGREE  OF MASTER  OF  S0IEI-JG3  III  QHEUI-STRY 


ln  Charge  of  Thesis 


A o Ting  Head  of  Department 


Recommendation  concurred  in* 


5 

5T 


Committee 


on 


Final  Examination* 


^Required  for  doctor’s  degree  but  not  for  master’s 


: - . . . : . 


TABLE  OF  CONTENTS 


page 


I.  Introduction  1 

II.  Review  of  Literature 4 

III.  Theoretical 8 

IV.  Experimental  10 

V.  Summary 18 

VI.  Bibliography 19 





ACKNOWLEDGMENT 

This  problem  was  brought  to  my  attention  in  January,  1922,  by 
Dr.  W.  S.  Putnam  of  the  Division  of  Industrial  Chemistry  at 
the  University  of  Illinois.  The  earlier  part  of  the  investiga- 
tion was  carried  on  under  his  direction.  The  latter  part  of 
the  work  was  done  under  the  supervision  of  Dr.  J.  H.  Reedy  of 
the  Inorganic  Division.  I wish  to  express  my  appreciation  for 
the  helpful  suggestions  offered  and  the  kindly  interest  shown 
by  both  men  on  all  occasions. 


THE  OXIDATION  OF  FERROUS  SULPHATE  TO  FERRIO  SULPHATE 

BY  MEANS  OF  AIR 

INTRODUCTION 

Among  the  wet  methods  for  the  extraction  of  copper  from  its 
ores,  leaching  with  ferric  sulphate  solution  is  among  the  oldest  if 
not  the  oldest  process. 

The  leaching  of  copper  ores  by  this  method  has  been  practiced 
at  Rio  Tinto , Spain,  since  1939  and  today  the  process  is  still  being 
applied  there  with  greater  success  than  at  any  other  place.  The 
process  there  is  called  cementation  and  the  product  cement  copper. 

Formerly,  the  ores  there  were  treated  by  ordinary  smelting 
methods,  but  the  sulphur  dioxide  liberated  destroyed  the  vegetation 
in  the  surrounding  country  and  became  such  a nuisance  that  further 
treatment  of  the  ores  by  the  then  used  process  was  prohibited  by  law. 

The  application  of  the  modern  method  was  the  result  of  inves- 
tigations conducted  by  FELIPE  PRIETO  of  Seville.  To  him  was  issued 
a royal  patent,  published  in  the  "Official  Gazette"  September  9,  1845 
described  as  "An  original  invention  for  a chemical  discovery  for 
utilizing  the  copper  bearing  pyritic  minerals  no  matter  how  low  the 
percentage  may  be". ( 1 ) 

The  method  consists  in  wetting  the  ore  heaps  with  water,  and 

with  proper  control  of  ventilation  of  the  pile,  the  mineral  is 
oxidized  by  air  and  moisture,  and  the  CJuSO^  which  is  formed  is 

washed  out 


-2- 


The  ore  is  principally  Cu2S  and  FeS2.  Advantage  is  taken  of 
the  following  reactions,  suggested  hy  Charles  H.  Jones.  (2)  (3) 

(1)  FeS2  + 70  + H20  — ? FeS04  + H2S04 

(2)  2FeS04  + H2S04  + 0 -* Fe2 f S04 )3  + H20 

(3)  Fe2(S04)3+  Cu2S  — > CuS04  + 2FeS04  + CuS 

(4)  Fe2 ( S04 }g  + CuS  + 30  + %0->CuS04  + 2FeS04  + H2S04 

It  is  seen  from  these  equations  that  the  Fe2(S04)g  is  the 
compound  which  reduces  the  copper  sulfides  to  a soluble  foxm.  Any 
copper  present  as  the  oxide  or  carbonate  would  also  undoubtedly  be 
dissolved  by  the  sulphuric  acid  formed  in  the  reactions  with  the 
sulfides. 

The  success  of  this  process  at  Rio  Tinto  naturally  led  to  oth- 
er attempts  to  leach  ores  with  ferric  sulphate  solutions.  Ores  whose 
copper  content  was  so  low  that  it  would  not  pay  to  treat  them  by 
ordinary  smelting  methods  seemed  especially  suited  to  this  treat- 
ment . 

Ferric  sulphate,  however,  is  a comparatively  expensive 
reagent  where  the  conditions  for  its  production  are  not  so  favorable 
as  at  Rio  Tinto.  Moreover,  it  is  reduced  to  ferrous  sulphate  in  the 
process.  It  therefore  became  necessary  to  find  some  cheap  means 
for  converting  this  ferrous  sulphate  back  to  ferric  sulphate  or  the 
expense  of  the  process  would  become  prohibitive. 

This  feature  is  unimportant  at  Rio  Tinto  because  the  ferric 
sulphate  used  is  foimed  in  the  ore  heap  as  is  shown  by  the  equations 
given  above.  The  liquors  are  simply  allowed  to  run  to  waste  in 
some  cases  after  the  copper  has  been  removed  by  precipitation  on 


scrap  iron. 


1 


1 

■4 

■ 

. 

. 

■ 


i 


j 


C ' ■ ' 


-Si- 


lt would  seem  that  a part  of  the  copper  might  he  lost  thru 
incomplete  removal  by  the  scrap  iron.  This  loss,  if  there  is  one, 
could  be  largely  eliminated  by  returning  the  waste  liquors  to  the 
ore  heaps  were  it  not  for  the  fact  that  the  basic  salts  of  iron, 
which  are  precipitated  when  solutions  of  ferrous  sulphate  are 
exposed  to  the  air,  have  a tendency  to  clog  the  ore  by  forming  a 
film  which  breaks  up  the  capillarity  of  the  ore,  preventing  access 
of  the  solvent  to  the  interior  of  the  piece.  Thomas  (4)  found  also 
that  the  presence  of  much  ferrous  sulphate  hindered  the  solvent 
action  of  ferric  sulphate  on  Cu2S. 

For  the  reasons  just  discussed,  it  became  imperative  to  find 
some  cheap,  fairly  efficient  and  fairly  rapid  my  to  oxidize  the 
ferrous  sulphate  liquors  to  ferric  sulphate.  If  this  could  not  be 
done,  the  ferric  sulphate  leaching  process  was  inapplicable  in 
places  where  the  character  of  the  ore  and  other  conditions  are  less 
favorable  than  at  Rio  Tinto. 

It  was  in  the  hope  that  some  light  might  be  thrown  on  this 
problem  of  the  regeneration  of  the  ferric- sulphate  liquors,  that 
this  investigation  was  undertaken. 


-4- 


RSVIEW  0?  THE  LITERATURE 

A good  many  articles  have  been  written  on  the  ferric  sulphate 
method,  and  several  processes  for  the  wet  extraction  of  copper  have 
been  patented,  which  are  based  on  the  regeneration  of  the  ferric 
sulphate  leach  liquors. 

Perhaps  the  most  interesting  and  instructive  article  is 
W.  L.  Austin's  (5)  (6)  account  of  experiments  carried  out  at  Cananea 
Mexico,  by  the  Cananea  Consolidated  Copper  Company. 

He  describes  the  process  by  which  the  oxidation  of  the  fer- 
rous sulphate  solutions  was  carried  out. 

A large  tank:  which  gave  a working  head  of  eleven  feet,  eight 
inches,  was  used  to  contain  the  solution.  The  belief  was  expressed 
that  a working  head  of  eighteen  feet  would  give  better  results. 

The  solution  was  heated  by  means  of  a steam  coil  and  preheated  air 
was  forced  in  from  the  bottom  agitating  the  solution  violently.  In 
one  run  which  was  kept  going  for  thirty  two  hours,  starting  with  a 
solution  containing  FeS04,  12.1$,  and  Pe^(S0^)g,  1.6$,  there  was 
present  at  the  end,  FeS04,  5.1$  and  Pe^fSO^Jg,  7.0$.  About  57.85$ 
of  the  original  FeS04  has  been  converted.  Trouble  was  experienced 
from  a basic  sulphate  which  formed  in  the  tank. 

In  regard  to  this  basic  sulphate,  Austin  says,  ,TThe  reactions 
which  take  place  when  an  attempt  is  made  to  oxidize  PeS04  to  the 
ferric  condition  without  the  presence  of  free  acid  are  very  compli- 
cated. Basic  ferric  salts,  of  which  there  are  many  varieties, 

invariably  form  and  are  precipitated,  thereby  causing  the  loss  of 
a large  part  of  the  iron  unless  free  H2S04  has  been  added  in  amounts 


-5- 


necessary  to  produce  the  neutral  Pe2( 304)5.  For  tiie  Purpose  speci- 
fied (the  formation  of  neutral  ferric  sulphate),  ten  parts  of  fer- 
rous sulphate  require  two  parts  of  concentrated  sulphuric  acid.  The 

reactions  which  occur  are  indicated  in  the  following  equation  : 

* 

2PeS04  + H2SO4  + 0 — ? Peg (304)5  + HgO 
"If  a solution  of  ferric  sulphate  is  heated  with  ferric  hydrate 
there  results  a deep  brown  liquor  containing  a more  basic  salt  - 
two  thirds  as  much  sulphuric  acid  combined  with  the  same  amount 
of  iron  as  in  the  neutral  salt.  This  basic  salt  is  also  formed 
when  a solution  of  ferrous  sulphate  is  slowly  oxidized  by  contact 
with  the  air,  while  at  the  same  time  a still  more  basic  salt  is 
produced  (with  one  sixth  as  much  sulphuric  acid  as  in  the  neutral 
sulphate)  together  with  other  soluble  sulphates  and  the  neutral 
ferric  sulphate. 

"The  basic  salt,  containing  two  thirds  as  much  sulphuric 
acid  a3  the  neutral  sulphate,  is  decomposed  by  heating  or  by 
dilution  of  the  solution,  the  resulting  products  being  neutral  fer- 
ric sulphate  and  a yellow  precipitate  containing  the  one  sixth  salt 
referred  to  above.  These  two  last  named  ferric  compounds  predomi- 
nate when  a solution  of  ferrous  sulphate  is  oxidized  by  exposure  to 
the  air,  and  are  claimed  by  some  authorities  to  be  the  final  products 
from  the  oxidation  described. 

10PeS04  + 50  — > 2Peg  ( SO4 ) 5 + Pe4S09" 

If  this  equation  be  representative  of  the  reaction,  it  is 

evident  that  40 % of  the  iron  and  \0%  of  the  acid  is  lost  ix  tne 
solution  is  oxidized  without  the  presence  of  free  acid. 

The  reaction  was  found  to  be  very  slow  and  the  proolem  of 


-6- 

oxidizing  the  ferrous  sulphate  solution,  it  was  thought,  was  the 

most  important  detail  of  the  successful  application  of  the  process. 

( 7}  v,  7) 

The  Irving  Leaching  Process:  - In  this  process,  the  FeS04 
solution  is  oxidized  hy  a steam  jet  which  agitates  it  violently, 
bringing  it  in  contact  with  the  air  and  supplying  heat.  The  basic 
sulphate  formed  is  dissolved  by  adding  H2SO4. 

In  an  article  entitled  "Some  Experiments  in  Heap  Leaching 
Copper  Ores",  (10),  Geo.  D.  Van  Arsdale  says,  "ITo  practicable  method 
of  cheaply  converting  for  leaching  purposes  ferrous  to  ferric  iron 
has  been  worked  out,  and  it  seemed  evident  that  the  only  sufficiently 
cheap  method  was  the  partial  conversion  to  be  obtained  by  evaporation 
Further,  if  this  evaporation  were  carried  on  in  intimate  contact  with 
the  ore  being  leached,  probably  a better  and  quicker  extraction 
could  be  obtained.  Accordingly,  the  following  steps  were  adopted 
as  a method  for  preliminary  tests:  (1)  Wetting  the  ore  with  excess  of 
solution  containing  ferrous  sulphate  and  a small  amount  of  ferric 
sulphate;  (2)  Allowing  the  ore  to  drain  and  air  dry  thoroly;  (3)  Pre- 
cipitating the  copper  from  the  resulting  liquor  by  iron,  returning 
the  liquor  to  ore  and  repeating. 

It  will  be  noted  that  this  method  is  identical  with  that 
practiced  at  Rio  Tinto,  except  that  the  ferrous  sulphate  is  supplied 
whereas  at  Rio  Tinto  this  in  unnecessary  because  of  the  nature  of 
the  ore,  the  iron  salts  being  formed  in  the  heap. 

Van  Arsdale  found  that,  using  the  method  outlined  above,  in 
which  a neutral  solution  of  ferrous  sulphate  was  used,  after  eight  or 
ten  leaching  cycles  the  solution  of  copper  practically  ceased.  The 
probable  reason  was  that  a yellowish  precipitate  of  ferric  hydroxide 

- _ 


-7- 


or  basic  sulphate  formed  and  broke  up  the  capillarity  of  the  ore, 
thus  preventing  the  solvent  action  of  the  solutions. 

In  a second  series  of  experiments,  a small  amount  of  acid  was 
added  to  the  lixivant  to  prevent  the  formation  of  this  precipitate. 
This  gave  better  results  and  the  recovery  of  copper  reached  a fixed 
rate  of  about  2$  of  the  copper  present,  for  each  leaching  cycle, 
instead  of  decreasing  to  zero  as  when  the  neutral  ferrous  sulphate 
was  used. 


-8- 


III.  THEORETICAL 

All  previous  experiments  on  the  oxidation  of  ferrous  sulphate 
solutions  with  air  seem  to  show  that  the  reaction  represented  "by 
the  equation  EEeSO^.  + HgSO^  + 0 — ^FegfSO^Jg  + HgO  is  a very  slow 
one  even  under  favorable  conditions,  and  that  the  energy  consumed 
(when  supplied  in  the  form  of  heat)  in  attempting  to  hasten  it  is 
very  great. 

If  we  calculate  the  heat  of  the  reaction  according  to  Hess’s 
Law,  from  the  heats  of  formation*  of  the  compounds  involved  in  the 
reaction  as  represented  by  the  equation  in  the  preceding  paragraph, 
it  will  be  seen  that  heat  evolved  by  the  reaction  is  equi valent  to 
176,400  calories  per  gram  mole  of  Peg( 804)3  formed.  This  heat  will 
of  course  be  dissipated  to  the  surroundings.  Then  according  to 
LeChatelier ' s principle,  the  reaction  once  started  should  proceed 
with  increasing  velocity  toward  the  right.  That  the  reaction  does 
start  and  does  proceed,  though  slowly,  toward  the  right  is  well 
known.  Why  does  it  not  proceed  rapidly  and  spontaneously  to  com- 
pletion? There  is  no  tendency  to  come  to  equilibrium  because  a 
solution  of  Peg ( SO4  )a>  has  no  tendency  to  reduce  except  in  the 
presence  of  strong  reducing  agents  such  as  nascent  hydrogen  and 
hydrogen  sulphide,  and  even  then,  care  must  be  used  if  complete 
reduction  is  desired.  On  the  other  hand,  PeS04  and  all  other  ferrous 
compounds  show  a tendency  to  pass  over  to  the  ferric  state.  They 
are  converted  q^iickly  and  completely  in  the  presence  of  strong 


*Heats  of  formation  taken  from  Landolt-Bornstein’ s TABELLSH. 


-9- 


oxidizing  agents.  Their  oxidation  is  also  easily  and  quickly 
accomplished  hy  electrolytic  methods. 

All  of  these  facts  would  indicate  that  the  oxidation  of 
FeS04  to  Peg (304)5  in  presence  of  air  and  free  acid  should  pro- 
ceed spontaneously  and  rapidly  to  completion.  As  a matter  of  fact, 
it  does  not.  We  have  a reaction  for  which  all  the  driving  forces 
are  favorable  and  yet  it  does  not  go.  This  being  the  conclusion, 
the  problem  seems  to  be  one  of,  first  finding  a suitable  catalyst, 
and  second,  making  the  mechanical  conditions  for  the  oxidation  as 
favorable  as  possible. 


-10- 


IV.  EXPSR IMEKTAL 

As  preliminary  experiments,  solutions  of  ferrous  sulphate 
were  aerated  hy  bubbling  air  through  them.  The  factors  varied 
were:  concentration  of  ferrous  sulphate,  concentration  of  free 

sulphuric  acid,  and  temperature.  The  temperatures  used  were 
room  temperature,  55°,  80°  and  95°C. 

These  tests  showed  that  the  concentration  of  ferrous  sul- 
phate within  the  limits  used  in  the  experiments  (2$  to  7$  by 
weight)  makes  little  difference  in  the  velocity  of  the  reaction; 
that  the  reaction  goes  somewhat  better  in  a neutral  solution  than 
in  an  acid  solution;  that  temperature  has  a marked  effect  on  the 
velocity  of  the  reaction,  increase  of  temperature  accelerating  it 
greatly.  The  period  of  these  experiments  was  twenty-four  hours, 
divided  into  two  periods  of  twelve  hours  each.  The  results  varied 
from  2.5$  oxidation  inthe  acid  solutions  which  were  aerated  at 
room  temperature,  to  30.9$  oxidation  in  the  neutral  solutions  which 
were  aerated  at  95°0. 

In  the  neutral  solutions,  basic  salts  of  iron  were  precipi- 
tated. These  salts  seemed  to  be  the  same  whether  formed  at  the 
higher  or  the  lower  temperatures.  Samples  were  analyzed  quantita- 
tively for  iron  and  sulphur,  the  method  of  analysis  being  as  follows: 

The  precipitates  were  filtered  onto  weighed  Gooch  filters 
and  dried  for  two  hours  at  110°.  The  residues  were  then  dissolved 
with  hot  concentrated  hydrochloric  acid.  The  iron  was  precipitated 
as  ferric  hydroxide  and  determined  as  Eeg03  by  the  ordinary  gravi- 
metric method.  The  filtrate  from  the  ferric  hydroxide  was  made 


-11- 


acid  and  the  sulphur  precipitated  and  determined  as  barium  sulphate. 

The  results  show  the  composition  of  the  samples  to  have  been 
Pe  68.95$,  SO4  9.71$.  This  leaves  21.54$  to  be  accounted  for 
as  oxygen  or  as  oxygen  and  hydrogen.  These  figures  represent  an 
average  of  four  determinations  which  were  consistent  within  0.4  of 
one  per  cent.  This  composition  would  indicate  a mixture  of  com- 
pounds, part  of  which  are  probably  hydroxides,  since  no  simple 
formula  would  fit  this  composition  unless  it  have  a very  large 
number  of  iron  and  oxygen  atoms  and  a small  number  of  SO4  groups. 
Presumably,  then,  it  may  be  considered  a mixture  of  basic  ferric 
sulphate  and  colloidal  ferric  hydroxide. 

On  ignition,  these  salts  decompose , leaving  as  a residue  a 
mixture  of  the  oxides  of  iron,  some  of  which  are  magnetic. 

Since  the  basic  salts  do  not  form  in  acid  solution  and  it  is 
obvious  that  H2SO4  in  some  form  must  be  supplied  if  the  normal 
Pegf 804)5  is  to  be  formed  from  Pe3C>4,  it  was  decided  to  try  to  use 
SOg  as  a basis  for  the  acid  at  the  same  time  the  solution  was  being 
oxidized. 

A mixture  of  S02  and  a large  excess  of  air  was  passed  thru 
a tube  furnace  filled  with  crushed  pumice  which  had  been  impregnated 
with  PegOg.  The  Pe^O^  was  expected  to  catalyze  the  union  of  a part 
of  the  oxygen  of  the  air  with  the  SOg  to  form  SO3.  The  pumice  was 
used  to  prevent  the  Pe203  from  packing  and  to  present  a large 
surface  of  the  catalyst  to  the  gas  mixture.  The  furnace  was  heated 
to  650°  0.  as  that  is  about  the  temperature  given  by  Partington( 11 ) 
as  the  most  favorable  for  the  catalytic  action  of  PegOg.  on  the 
union  of  SO2  and  oxygen.  The  hot  gases  from  the  furnace  were 
discharged  directly  into  a solution  of  ferrous  sulphate  which  was 


-12- 

stirred  vigorously  with  a motor. 

In  using  this  apparatus,  it  was  found  that  unless  the  ratio 
of  air  to  SC2  was  kept  very  large,  a part  of  the  SOg  came  through 
into  the  solution  and,  since  it  is  a reducing  agent,  almost  entirely 
prevented  the  oxidation  of  JeSC^  to  FegfSO^.  )g.  When,  however,  a 
very  small  amount  of  SO2  was  mixed  with  a relatively  large  amount 
of  air,  the  oxidation  went  forward  slowly. 

In  one  10-hour  run  which  may  "be  taken  as  representative,  the 
solution,  containing  at  the  start  of  the  run  5 fo  by  weight  of  FeSO^. 
and  only  a trace  of  FepfSO^)^  , was  subjected  to  the  action  of  the 
hot  gases  as  they  were  dischargedirom  the  furnace.  Yfater  was  added 
from  time  to  time  to  make  up  for  losses  by  evaporation.  A thermome- 
ter in  the  solution  showed  an  average  temperature  of  70°.  SampJ.es 
were  taken  at  intervals  of  two  hours  and  analyzed  for  iron  present 
as  Fe++  and  Fe+++.  Ho  basic  salts  were  precipitated. 

The  potassium  permanganate  method  was  used  in  these  and  all 
other  analyses  made  for  ferrous  and  ferric  iron  in  connection  with 
these  experiments.  Ten  cc.  of  the  iron  sulphate  solution  were 
titrated  against  potassium  permanganate  solution,  giving  the  iron 
present  as  Fe++ . A second  ten  cc.  of  the  iron  sulphate  solution 
were  run  through  a Jones  reductor  and  then  titrated  against  the 
permanganate,  giving  the  total  iron  present.  The  difference  of 
course  gives  that  part  of  the  iron  present  in  the  ferric  state. 

The  results,  which  were  very  disappointing  were  as  follows: 


f 


-13- 


TIMS 

fo  ¥E++ 

fo  FE+ 

0 hrs. 

99.5 

0.5 

2 n 

97.2 

2.8 

4 IT 

94.8 

5.2 

6 " 

93.9 

6.1 

8 " 

93 . 7 

6.3 

10  " 

93.4 

6.6 

There  was  appreciable  oxidation  during  the  first  six  hours,  but 
after  that,  the  reaction  though  continuous  was  very  slow.  The 
results  on  the  whole  were  less  satisfactory  than  when  air  alone 
was  used.  The  only  gain  over  air  was  that  acid  did  not  have  to  be 
added  and  this  was  more  than  offset  by  the  heat  used  in  keeping  the 
temperature  of  the  furnace  up  to  650°. 

It  became  evident  that  heat  was  inefficient  as  a catalyst 
and  that  the  S02  method  here  used  for  supplying  the  needed  sulphuric 
acid  was  unsatisfactory. 

Various  substances  were  successively  tried  as  catalysts 
with  uniformly  poor  results. 

Animal  charcoal  was  suspended  in  the  solution  and  the  whole 
stirred  vigorously  with  a motor  while  air  u©s  passed  through. 

This  experiment  was  tried  both  with  and  without  heating  the  solu- 
tion. Precipitated  silica  and  manganese  dioxide  were  tried  under 
the  same  conditions  as  the  charcoal.  There  seemed  to  be  no 
tendency  for  any  of  these  to  catalyze  the  reaction. 

It  was  next  decided  to  try  the  effect  of  the  oxides  ox  some 
of  the  heavy  metals  which  have  two  or  more  oxides  as  it  is  known 

that  some  of  these  oxides  will  catalyze  certain  reac t ions  ♦ The 


» 


. 


. 


-14- 


metals  which  meet  this  requirement  include  chromium,  cobalt,  lead, 
manganese,  and  mercury  among  the  common  metals.  Others  such  as 
platinum  are  unsatisfactory  on  account  of  the  expense  if  for  no  othei 
reason.  Of  the  oxides  of  the  more  common  metals  mentioned,  it 
happens  that  manganese  dioxide  is  the  only  one  that  is  not  soluble 
in  dilute  solutions  of  sulphuric  acid. 

At  this  point,  two  pieces  of  apparatus  were  devised  with  the 
idea  of  improving  the  mechanical  conditions  for  the  oxidation  by 
providing  greater  surface  contact  between  the  solution  and  the  air 
and  between  the  solution  and  the  catalyst. 

to 

In  the  first,  the  solution  was  made/\flow  over  the  surface  of 
a large  clay  crucible  (area  of  surface  about  420  square  centimeters) 
at  the  rate  of  one  hundred  cubic  centimeters  per  minute.  No  heat  or 
catalyst  was  used  and  the  results  were  unsatisfactory. 

In  ths  second  apparatus,  a diagram  of  which  appears  on  the 
next  page,  the  solution  was  caused  to  trickle  down  through  a tower, 
packed  loosely  with  the  catalytic  material,  while  air  was  blown  into 
the  tower  from  the  bottom.  The  solution,  after  passing  over  the 
catalyst  and  at  the  same  time  being  aerated,  drained  back  into  the 
reservoir  and  was  returned  to  the  top  of  the  tower  again  by  a lift 
pump  operated  by  a current  of  air.  The  cycle  could  be  made  contin- 
uous as  long  as  desired. 

As  packing  for  the  tower  crushed  pumice  which  had  been  impreg- 
nated with  manganese  dioxide  was  used.  The  impregnation  of  the 
pumice  was  accomplished  by  boiling  with  a concentrated  solution  of 
EMn04  followed  by  heating  to  red  heat  in  a muffle  furnace  for  one 

hour,  or  reducing  the  EMnQ^  with  alcohol  and  drying. 


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-OXIDISING  APPARATUS- 

A.  Catalyst  tower 

B.  Lift  pump 

C.  Reservoir  for  solution 

D.  Return  drain 


-16- 


A run  was  made  with  this  material  as  packing  for  the  tower. 

The  particles  of  manganese  dioxide  washed  off  easily  and  became 
coated  with  the  basic  salts  described  above,  but  the  results  were 
sufficiently  encouraging  to  warrant  further  investigation  along 
this  line. 

In  the  next  experiment,  the  tower  was  packed  with  lumps  of 
mineral  manganese  dioxide  (pyrolusitet  graded  to  about  one  fourth 
inch  sizes.  The  height  of  the  column  of  catalyst  was  eighteen 
inches.  The  solution  used  was  one  normal  with  respect  to  FeSO^ 
and  contained  the  theoretical  amount  of  H2SO4  necessary  to  form 
Peg (SO^Jg  . Air  was  blown  into  the  bottom  of  the  tower  as  rapidly 
as  was  possible  without  blowing  the  solution  and  the  catalyst-  out  of 
the  top.  The  solution  was  run  into  the  top  of  the  tov/er  at  the 
rate  of  one  hundred  cubic  centimeters  per  minute. 

After  all  of  the  solution  had  passed  through  the  tov/er  once, 
(the  capacity  of  the  apparatus  was  approximately  two  liters)  a 
sample  was  taken  for  analysis.  The  results  exceeded  all  expecta- 
tions. It  was  found  that  eighty- two  per  cent  of  the  iron  had  been 
oxidized  to  the  ferric  state.  By  the  time  the  solution  had  passed 
five  times  over  the  catalyst,  the  oxidation  was  practically  one 
hundred  per  cent  complete. 

This  result  has  been  repeated,  without  changing  the  catalyst, 
using  solutions  of  varying  FeS04. concentration.  The  catalyst  becomes 
coated  with  the  basic  sulphates,  especially  where  the  pieces  touch 

each  other  or  the  walls  of  the  tower.  It  is,  hov/ever,  easily 
cleaned  without  removing  from  the  tower  by  washing  with  a little 
dilute  sulphuric  acid.  Where  solutions  are  used  to  which  the 


. 


-17- 


theoretioal  amount  of  acid  has  been  added  for  the  formation  of 
normal  ferric  sulphate,  it  is  a long  time  before  the  lumps  of 
manganese  dioxide  become  sufficiently  coated  to  interfere  seriously 
with  their  catalytic  efficiency.  It  is  thought  that  if  larger 
lumps  were  used,  the  tendency  to  become  coated  would  be  less  because 
as  stated  before,  the  coating  formed  more  noticeably  where  the 
pieces  were  in  contact  with  each  other. 

The  same  manganese  dioxide  was  used  for  all  runs  and  during 
these  experiments  showed  no  tendency  to  become  ,Tpoisoned,T  or  lose 
its  efficiency  except  on  account  of  the  coating  described  in  the 
preceding  paragraph. 


-18- 


v.  summary 

A process  has  been  developed,  for  the  rapid  oxidation  of 
ferrous  sulphate  in  solution  to  ferric  sulphate  by  means  of  air. 

ITo  heat  or  other  form  of  energy  is  required  except  that  necessary 
for  circulating  the  solution  and  blowing  the  air  into  the  tower. 

The  apparatus  is  simple  and  capable  of  modification  to  suit 
large  scale  operations.  By  increasing  the  height  of  the  catalyst 
column  or  by  connecting  a number  of  towers  in  series,  it  should  be 
possible  to  ma&e  the  process  continuous  and  at  the  same  time  to 
secure  practically  one  hundred  per  cent  oxidation. 

The  catalyst  in  view  of  the  amount  required  is  comparatively 
inexpensive.  The  manganese  ore  has  been  quoted  on  the  market 
within  the  year  at  fifty-five  dollars  a ton. 


_J  i 


• 

V. 


1 


. • r 


-19- 


VI.  BIBLIOGRAPHY 

1.  Mining  and  Scientific  Press,  Vol.  122,  p.  185. 

2.  Transactions  of  the  American  Institute  of  Mining  Engineers, 
February,  1904,  Chas.  H.  Jones. 

3.  Mineral  Industries,  Vol.  XII,  p.  112,  1903. 

4.  Metallurgie,  1904,  p.  62. 

5.  Mines  and  Methods,  September,  1910,  W.  L.  Austin. 

6.  Hydrometallurgy  of  Copper  (Greenwalt)  p.  194,  et  seq.. 

7.  Mining  and  Scientific  Press,  Vol.  108,  p.  77. 

8.  Metallurgical  and  Ohemical  Engineering,  Vol.  XI,  p.  160. 

9.  U.S.  Patent  1,048,541,  December  31,  1912. 

10.  Engineering  and  Mining  Journal,  Vol.  105,  p.  225. 

11.  TextbooX  of  Inorganic  Chemistry  (Partington) 


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